1—Spectrophotometric Analysis of
Commercial Aspirin
Name: ___________________________________________________
Date: _____________________________________________________
Section: _________________________________________________
Objectives
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Learn about the absorption of light by molecules
Learn the basic components of a spectrophotometer
Learn to prepare standards diluting a stock solution
Prepare a Beer’s Law curve from standard solutions
Use a Beer’s Law curve to calculate the concentration of an unknown substance
Gain experience pipetting, a technique you learned in CHEM 131L
Gain experience weighing small samples
Pre-Laboratory Requirements
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Read Chapter 7.1 – 7.3 in Silberberg
Watch the instructional videos titled “Mass Balance/Analytical Balance” and “Pipetting”
Pre-Lab Questions (if required by your instructor)
Laboratory Notebook—prepared before lab (if required by your instructor)
Safety Notes
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Eye protection must be worn at all times
Sodium hydroxide is caustic and should not come in contact with your skin or clothing. Wear gloves
when handling this chemical. A lab coat or lab apron is recommended.
Discussion
Light is a form of electromagnetic radiation. We are most familiar with the visible portion of the electromagnetic
spectrum because this is the region of light to which our eyes are sensitive. Visible light, however, is only a small
segment of the entire electromagnetic spectrum (see Silberberg, pg. 218).
Electromagnetic radiation is a form of energy that may be represented as a wave or as a particle. The wave model
for light is more useful for predicting the behavior of light in our day to day activities, but at an atomic scale light
is better described by particles called photons.
The wave model for light describes electromagnetic energy in terms of wavelength, frequency and intensity (see
Figure 1). One wavelength is represented by the time from one peak to the next in any wave front (also, one
trough to the next). As the wavelength of the radiation becomes shorter, a larger number of waves per unit of time
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(i.e., the frequency), becomes greater. Wavelength is often expressed in meters, or nanometers for visible light,
and frequency is usually expressed in hertz (Hz). One Hz is one cycle per second. The energy of light is calculated
from frequency using the Planck equation:
E = hʋ
where E is the energy in joules, h is Planck’ constant (6.63 x 10-24 J·s), and ʋ is the frequency in Hz.
Figure1. The wave model for light
The symbol for wavelength is the Greek letter lambda, λ, and the symbol for frequency is the Greek letter nu, ʋ.
The speed of light in a vacuum is 3.0 x 108 m/s. Since the speed of light is a constant, we can use it to calculate
wavelength if we know the frequency of radiation (see sample Problem 7.1 in Silberberg). The relationship is:
λʋ = c
where λ is the wavelength of radiation, ʋ is the frequency of radiation, and c is the speed of light (3.0 x 108 m/s).
White light is composed of all the wavelengths within the visible region of the spectrum (see Figure 7.3 in
Silberberg). When white light falls on an object, some of the incoming radiation may be absorbed. Wavelengths
that are not adsorbed will be transmitted from the object. In the example below, the white egg does not adsorbed
any of the incoming radiation, and therefore appears white. The apple adsorbs blue and green radiation from white
light, transmitting red light. Therefore, the apple appears red.
Figure 2. The white egg in this illustration does not absorb red, green or blue light, and appears white when
illuminated with white light. The apple absorbs green and blue light, and transmits red light. The apple therefore
appears red.
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Colors are often used to identify objects. The spectrum of an object is a graph of the absorbance of that object
plotted against the wavelength of light. At the molecular level, spectra are used to identify unknown molecules.
The peaks in the spectrum can be used to identify unique components within a molecular structure. Figure 7.11 in
your textbook is a spectrum for chlorophyll. Notice that the chlorophyll spectrum has intense adsorption peaks in
the blue and the red regions of the spectrum. Why do you think chlorophyll is green?
Spectrophotometers disperse white light into its component wavelengths by passing light through either a prism or
a grating. The intensity of light at any wavelength can be measured with a photocell, a photomultiplier tube, or a
solid-state device called a charge-coupled device (CCD). The optical path of the Spectronic 20 used in today’s
experiment is shown below:
Figure 3. Optical path of the Spectronic 20 spectrophotometer (Courtesy of Fisher Scientific).
Absorbance is the ratio of the negative logarithm of light intensity transmitted from a sample divided by the
intensity of incoming light. This is expressed mathematically as:
−𝑙𝑙𝑙10
𝐼
=𝐴
𝐼𝑜
This relationship can be used to measure the amount of material present by measuring the intensity of the
adsorption peak at a specific wavelength. In today’s experiment we will measure the intensity of the colored
complex that forms when iron (III) is mixed with aspirin to determine the amount of pure aspirin (acetylsalicylic
acid) in commercial aspirin tablets.
The Beer-Lambert Law (often shortened to Beer’s Law) relates the absorbance of a sample to the concentration of
a species in solution and is the relationship used when making quantitative measurements. Mathematically, Beer’s
law is expressed as:
𝐴 = 𝝐𝑙𝑙
where A is the measured absorbance of the solution, 𝝐 is the molar absorptivity of the substance, l is the path
width for the cell, and c is the concentration. By measuring the absorbance of several solutions of known
concentration we are able to prepare a graph that can be used to determine concentrations of unknowns.
A graph of absorbance vs concentration is called a Beer’s Law curve in honor of the chemist who first discovered
the relationship between absorbance and concentration. Figure 4 is Beer’s Law curve for the absorbance of an
iron-salicylate complex (the substance prepared in today’s experiment) plotted against different concentrations.
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Figure 4. Beer’s Law plot of Fe(III)-salicylate complex.
Procedure
Part I. Preparation of Standards and of the Aspirin solutions
1. Weigh out approximately 0.16 g of acetylsalicylic acid and record the exact mass on your report sheet or in
your notebook.
2. Transfer this sample into a 125 mL Erlenmeyer flask.
3. Add 5 mL of 1 M sodium hydroxide and carefully heat the mixture until all solid dissolves.
4. Allow this solution to cool, and then transfer it into a 100.0 mL volumetric flask with the use of a glass
funnel. Rinse the flask with DI water to insure complete transfer.
5. Dilute the solution with deionized water to the 100.0 mL mark on the flask (Note: Be sure to label this
solution as your stock solution). Invert the flask several times to insure the sample is thoroughly mixed.
6. Using a 1 mL graduated pipet, transfer a 0.5 mL sample of this stock solution into a 10.0 mL volumetric flask
and dilute this solution to the 10.0 mL mark with 0.02 M iron(III) chloride.
7. Place this solution in a test tube labeled solution A.
8. In a similar fashion, prepare solutions labeled B, C, D, and E by using 0.40, 0.30, 0.20, and 0.10 mL aliquots
of the sodium salicylate solution, diluting to 10.0 mL with iron(III) chloride solution.
9. Obtain one commercial aspirin tablet and break it into two approximately equal pieces. Record the exact mass
of each piece on your report sheet or in your notebook.
10. Transfer the two aspirin pieces into two 125 mL Erlenmeyer flasks, and label the flasks Sample 1 and Sample
2.
11. Add 5 mL of 1 M sodium hydroxide and carefully heat the nixture until all solid dissolves.
12. Allow these solutions to cool, and then transfer then into two 100.0 mL volumetric flasks, using a glass funnel
to insure a quantitative transfer.
13. Dilute these solutions to the 100.0 mL mark on the flask, and label these flasks Sample 1 and Sample 2.
14. Invert the volumetric flasks several times to insure the samples are thoroughly mixed.
15. Using a 1 mL graduated pipet, transfer a 0.3 mL sample of each solution into two 10.0 mL volumetric flasks
and dilute to the 10.0 mL mark with 0.02 M iron(III) chloride. Label these flasks Sample 1 and Sample 2.
16. Transfer solutions to test tubes, and label the test tubes Sample 1 and Sample 2.
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Data
Mass of acetylsalicylic acid, (±0.001 g)
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Moles of acetylsalicylic acid (mol)
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Concentration of acetylic acid in 100.0 ml volumetric flask (Stock solution)
__________
Mass of aspirin sample 1 (±0.001 g)
__________
Mass of aspirin sample 2 (±0.001 g)
__________
Part II. Measure Absorbance of Standards and Aspirin Samples
1. With the assistance of your instructor, adjust the 0% Transmittance and 100% Transmittance settings on the
Spectronic 20. Refer to the Operating Instructions at the end of this experiment for guidance. Note that you
must use the iron(III) chloride solution to set the zero absorbance (100% T).
2. Record the absorbance for standard solutions A, B, C, D and E.
3. Record the absorbance for Sample 1 and Sample 2.
Solution
Concentration
Absorbance
A
__________
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B
__________
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C
__________
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D
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E
__________
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Sample 1
__________
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Sample 2
__________
__________
Calculations
For this example we assume 0.400 g of acetylsalicylic acid (aspirin) is treated as outlined in the procedure (you
should use the actual mass recorded on your report sheet). The concentration of complex in the stock solution can
be found as follows:
The molar mass of acetylsalicylic acid (C9H8O4) = 180.2 g/mol
0.400𝑔 𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎 𝑎𝑎𝑎𝑎 𝑥
1 𝑚𝑚𝑚
180.2 𝑔
The concentration of the acetylsalicylic acid stock solution then is:
= 2.22𝑥10−3 mol C9H8O4
2.22x10-3 mol C9H8O4 ÷ 0.1000 L = 2..22x10-2 M
This stock solution is used to prepare standards for the Beer’s Law curve by diluting aliquots of the stock solution
into 10.00 mL volumetric flasks, and diluting to volume with iron(III) chloride solution. The concentration of
C9H8O4 in the standards is calculated with the relationship M1V1 = M2V2, where M1 if the concentration of the
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stock solution, V1 is the volume of the stock solution transferred, V2 is the volume of the diluted solution (10.00
mL for all standards in this experiment), and M2 is the new concentration.
For example, the concentration of C9H8O4 in the solution A would be calculated as follows:
𝑀1 𝑉1 = 𝑀2 𝑉2 = 2.22𝑥10−2 𝑀𝑀0.50𝑚𝑚 = 𝑀2 𝑥10.00𝑚𝑚
𝑀2 =
0.50 𝑚𝑚 𝑥 2.22𝑥10−2 𝑀
= 1.11𝑥10−3 𝑀
10.00 𝑚𝑚
You can plot your absorbance data for standards A, B, C, D and E on the graph paper included with this
experiment, or you can plot the data with Excel as an X-Y scatter plot. The concentration data should be plotted
on the x-axis and the absorbance data on the y-axis. If you use Excel to plot the data, add a trend line after the
graph has been generated. Your graph should look like the one in Figure 4.
Use the graph you just prepared to find the concentration of aspirin in Samples 1 and 2. Use the trend line to find
the concentration that would give the absorbance recorded for Sample 1, and write this number on the data sheet.
Find the concentration for Sample 2, using the same procedure, and write this number on the data sheet. You can
also use the equation from the trend line to find concentrations by entering you absorbance value for Y in the
equation, and solving for X.
The concentration obtained from the Beer’s Law plot represents the concentration of acetylsalicylic acid in the
10.0 mL volumetric flask. To find the mass of acetylsalicylic acid in your original sample you must correct for the
dilutions that were done preparing the sample. This is simply the reverse of the process just described.
Calculation of percent aspirin in commercial aspirin tablets
Sample 1
Sample 2
Mass of sample
__________
__________
Absorbance from graph
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__________
Concentration in diluted solution
__________
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Concentration in Sample Stock solution
__________
__________
Moles of acetylsalicylic acid in Stock solution
___________
__________
Mass of acetylsalicylic acid in Stock solution (g)
___________
__________
Percent acetylsalicylic acid in sample
___________
__________
Average percent acetylsalicylic acid
___________
Show sample calculations here:
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Operating Instructions for the Spectronic 20
Instrument Start Up: Your TA should have already turned on the instruments so they warm up during the short
lecture before the experiment begins. If he/she has not, turn on the instrument by rotating the left-hand knob, the
amplifier control, clockwise. On some models of instruments, a red jewel will light at this point. Allow about 10
to 15 minutes for the instrument to warm up before recording any measurements.
Setting 0% Transmittance: After the instrument has warmed up, adjust the left-hand knob so that the meter needle
reads 0 on the percent transmittance scale. This adjustment is made with no cell in the holder, so that no light
strikes the phototube.
Cleaning Cuvettes and Obtaining a Matched Pair: Rinse two test tube cells, called cuvettes, with deionized water.
Fill each about two-thirds full with water and insert one cuvette into the sample holder. Adjust the right-hand
knob so that the needle reads about 50% transmittance. Replace this cuvette with the other. If the %T reading of
the second cuvette is not within 1% of the value read for the first cuvette, repeat using other cuvette combinations
(Note: Do not handle the lower portion of the cuvettes. Wipe off the outside of the cuvettes with Kimwipes before
each reading).
Setting 100% Transmittance: Adjust the wavelength to the desired setting by using the knob on the top of the
instrument. Turn the right-hand knob, the light control, counterclockwise almost to its limit before inserting a
cuvette into the sample holder. Insert the cuvette containing the blank into the holder, matching up the index line
on the cuvette exactly with the index line on the holder. Close the top of the holder. Adjust the right-hand knob
clockwise until the needle reads 100 on the percent transmission scale. Immediately remove the test tube to avoid
fatiguing the phototube.
Check of 0% T and 100% T: After the cell is removed from the holder, an occluder automatically drops into the
light beam. The needle should then read 0. Each time the wavelength is changed and during all measurements at
the same wavelength, stablish or check the 0% T and 100% T settings.
Sample Measurements: Insert your sample and close the lid on the sample compartment. Record the percent
transmittance on the meter for each sample.
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Absorbance
Concentration
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